AP Chemistry Name __________________________ Ms. Ye Date ________________ Block ____

Do Now:

1. Complete the table based on the example given

Location Element Electron Configuration Metal, Nonmetal or Semi-metal Metalloid)?

Group/Family Name

Group 1,

Period 1

Hydrogen

(H)

1s1 Nonmetal (none)

Group 11,

period 4

Group 14,

Period 3

Group 17,

Period 4

2. Why do all elements want to obtain a noble gas electron configuration?

3. Which of the following elements has the most similar properties to Ca?

(elements in the same ________________ have the most similar properties because

they have the same ________________________________________________________)

a. K b. Sc c. Sr d. Ar

Review-Properties of Elements & Periodic Trends:

Effective Nuclear Charge-a measure of the positive attractive force of the nucleus towards negatively charged electrons due to the number of _______________________________; how much attractive force an electron feels can be affected by the number of shielding electrons (defined below)

Electron Shielding Effect-electrons in the energy levels closer to the nucleus protects the electrons in the outer shells and lessens the effect of the positive, attractive force of the nucleus. _________________________________________________________________

1. Atomic Radius: size of an atom

When looking at elements going down a GROUP, atomic radius increases

o As you go down a group, more _____________________________________________

are being added, therefore increasing the size.

When looking at elements going across a PERIOD, atomic radius decreases

o As you go across a period, the ________________________________________

therefore the nucleus _______________________________________________ the

electrons of the atom, and the radius decreases

Examples: For each pair of elements below, circle the one with the larger atomic radius.

a. Na and Cl

b. Mg and Sr

c. C and B

d. Ar and Ne

e. K and Se

f. Sb and B

g. Br and Ca

h. Ge and C

**Use nuclear charge to explain/show why Be has a smaller atomic radius than Li. Include a

Bohr diagram for both Li and Be.

2. Electronegativity: tendency to attract or gain an electron

When looking at elements going across a PERIOD, electronegativity increases

o As you go across a period, the _____________________________________________

therefore the nucleus ____________________ (gain) ___________________________

o Elements towards the right side of the periodic table are “closer to becoming a noble

gas—they want to “gain electrons”

o Exception: _____________________________

When looking at elements going down a GROUP, electronegativity decreases

o As you go down a group, the atomic radius increases. The _______________________

______________________________________________________________, therefore

_____________________________________________ and the attraction for electrons

Examples: For each pair of elements below, circle the one with the greater Electronegativity.

a. Na and Cl

b. Mg and Sr

c. C and B

d. Ar and Ne

e. K and Se

f. Sb and B

g. Br and Ca

h. Ge and C

3. Ionization Energy: energy required to remove an electron

When looking at elements going down a GROUP, ionization energy decreases

o As you go down a group, the atomic radius ______________________. With more shells

being added, the_________________________________________________________,

therefore _____________________________________________ and attractive force on

the electrons, making them easier to remove.

When looking at elements going across a PERIOD, ionization energy increases

o As you go across a period, the ___________________________________________

therefore the ______________________________________________________________

making it harder to remove an electron

o Elements towards the right side of the periodic table “don’t want to lose electrons” (they

want to gain electrons) to become like a noble gas. Therefore, it is difficult (requires more

energy) to remove an electron

Examples: For each pair of elements below, circle the one with the greater Ionization Energy.

a. Na and Cl

b. Mg and Sr

c. C and B

d. Ar and Ne

e. K and Se

f. Sb and B

g. Br and Ca

h. Ge and C

**Use electron shielding to explain why Mg has a lower ionization energy and a lower

electronegativity than Be. Include a Bohr diagram for both Mg and Be.

More about Ionization Energy….

First Ionization Energy –energy required to remove 1 electron

Second Ionization Energy – energy required to remove 2nd electron

Third Ionization Energy – energy required to remove 3rd electron

In general 1st I.E. ______ 2nd I.E. ______ 3rd I.E. Because as the # of electrons decreases, the

nucleus (# protons doesn’t change) has a stronger pull on the electrons that are remaining.

Based on the relative “jump” between ionization energies, you can tell how many

___________________________ the element has

1.

a. Between which 2 ionization energies do you see the biggest jump?

b. How many valence electrons would this element have?

c. What group would this element be found in?

2. What group would this element be found in?

Ions and Ionic Radius When an atom loses electrons and becomes a cation, its radius becomes _______________ than that of

the neutral atom

o # protons _______________________________# electrons, therefore increasing the effective nuclear

charge, meaning that there is a stronger pull of the electrons towards the nucleus.

When an atom gains electrons and becomes an anion, its radius becomes _______________ than that of a

neutral atom

o When electrons get added to the same energy level, they repel each other

*Note: the term isoelectronic refers to ____________________________________________________

Properties of Metals vs. Nonmetals

Metals Nonmetals

Malleable (can be hammered/molded into sheets)

Ductile (can be drawn/pulled into a wire)

Have luster (are shiny when polished)

Good conductors (allow heat & electricity to flow throw them)

Not malleable or ductile; instead, they

are brittle (shatter easily)

Lack luster; instead, they are dull

They are either poor or nonconductors

Reactivity of Metals vs. Nonmetals

Reactivity of a metal is related to its __________________________________________

o The _______________________________________________________________,

the _______________________________________________________ the metal

o Trend (within the metals on the periodic table):

Going down a group: __________________________________________

Going across a period: _________________________________________

o Most reactive metal: _______________________________________________

Reactivity of a nonmetal is related to its __________________________________________

o The____________________________________________________________________,

the _________________________________________________________ the nonmetal

o Trend (within the nonmetals on the periodic table):

Going down a group: __________________________________________

Going across a period: _________________________________________

o Most reactive nonmetal: ___________________________________________________

Periodic Trends Questions: 1. Which of the following atoms has the smallest atomic radius?

a. Li *Explain your answer using the term effective nuclear charge: b. Be c. C d. F

2. As the elements of a group are considered from top to bottom, the atomic radius a. Increases b. Decreases

3. Which element in group 17 is least likely to lose an electron? a. Chlorine *Explain your answer using ionization energy and nuclear charge b. Iodine c. Bromine d. Fluorine

4. Of all the elements, the one with the highest electronegativity is found in period a. 1 *What is the identity of this element? How do you know? b. 2 c. 3 d. 4

5. As the elements in group 2 are considered in order of increasing atomic number, what happens to the atomic radius? Why?

6. Fill out the following table about metals and nonmetals Metals Nonmetals

Location on Periodic Table

Lose or Gain electrons to obtain noble gas electron configuration?

Form cations or anions?

Relative ionization energy (high or low)

Relative electronegativity (high or low)

*Explain how the radius, along with electron shielding, would affect the ionization energy as you consider elements going down the group:

Model 1: Hydrogen and Lithium

Hydrogen Atom Lithium Atom

1. In the Hydrogen and Lithium atoms, what force of attraction holds the electron(s) in the atom?

2. The amount of energy necessary to remove an electron from an atom is called the ionization energy of that electron. What is the relationship between the ionization energy of an electron and the net attractive force that holds an electron in an atom?

3. Consider the electrons in an atom of lithium as diagrammed in the model above. Which electron, 1 or 3, will require more energy to be removed? Support your answer by discussing the attractive forces in the atom and how they might be different for electrons 1 and 3

Model 2: Photoelectron Spectra of Lithium

Refer to the PES graph for Lithium above: 4. What are the units of the x-axis? What is unusual about the way the x-axis values are

graphed?

5. Which of the peaks in the graph represents electrons that are more tightly held by the nucleus? Explain your reasoning

6. The number of peaks in a PES spectrum reveal the number of energy sublevels occupied by electrons in an atom. Based on the energy values of the peaks, label each peak with the electrons in a lithium atom to which they correspond.

7. Why is the higher energy peak about twice as high as the lower energy peak?

8. Using the lithium PES spectrum as a starting point, sketch how the spectrum of the next larger element (Beryllium) would look like. Recall that Be will have one more proton in its nucleus and on more electron in its sublevels.

Model 3: Bohr Model Modified—Photoelectric Spectra of Neon

Bohr Model of Ne

Shell model of Ne PES of Ne

9. Based on the PES of Ne, which is a better electron model – Bohr or Shell model? Why? 10. In the table below, label the peaks for the PES of Ne with the correct sublevels and number

of electrons.

11. Write the electron configuration for Neon. 12. What do you notice about your answers to 10 and 11?

Peak A B C

Sublevel

# of electrons

A B

C

SUMMARY: 13. What does the PES tell us about the electrons in an atom? How can you use the PES of an

atom to determine the identity of the atom? Practice #1

14. Label the peaks with the correct sublevels and number of electrons for each atom.

15. Consider the attractive and repulsive forces in the atoms of sulfur and phosphorus.

a. Explain why most of the peaks in the sulfur spectrum are shifted to the left relative to the

peaks in the phosphorus spectrum.

b. Explain why peak E in the sulfur spectrum is shifted slightly right compared to peak E in the

phosphorus spectrum.

16. Sketch the PES spectra for chlorine using the spectra for P and S as a guide

Practice #2 17. Answer the following questions based on the principles of atomic and electronic structure.

Element #1

Element #2

a. The diagrams above represent the photoelectron spectra (PES) for two different elements located in Period 3. Identify Element #1 and Element #2. Label each peak in each PES with the name of the orbital in which the electrons are located. Identify the orbital with its principal quantum number (n) and orbital type, such as 1s or 3p.

Identity of

Element #1 A B C D

Identity of

Element #2 A B C D E

b. Explain why peak E is twice as high as peak D in the PES for Element #2. c. Explain why peak A in Element #2 has a higher binding energy than peak A in Element #1.

Binding Energy (eV)

Rel

ativ

e N

um

ber

Of

Ele

ctro

ns

10,000 1000 10 1100

Binding Energy (eV)

Rel

ativ

e N

um

ber

Of

Ele

ctro

ns

10,000 1000 10 1100

A B

C

D

A B

C

D

E

Practice #3

18. Look at the PES below. Identify the sublevel & number for each peak; write the electron

configuration; and identify the element

19. Look at the PES below. Identify the sublevel & number for each peak; write the electron

configuration; and identify the element

Electron configuration:

Element identity is

Electron configuration:

Element identity is